chemical bonding and bond formations in molecules

chemistry

1. CHEMICAL BONDING

2. Chemical Bond It is defined as the attractive forces which hold the various chemical constituents (atoms, ions, etc.) together in different chemical species. Kossel-Lewis Approach to Chemical Bonding According to this theory. atoms take part in the bond formation to complete their octet or to acquire the electronic configuration of the nearest inert gas atoms (Octet rule). This can be achieved by gaining, losing or sharing the electrons.

3. Covalent Bonds : These occur when two atoms share one or more pairs of electrons. The shared electrons allow each atom to attain the electron configuration of a noble gas. Covalent bonds can be single, double, or triple, depending on the number of electron pairs shared. Ionic Bonds : These form when one atom transfers one or more electrons to another atom, resulting in the formation of positively and negatively charged ions. The electrostatic attraction between these oppositely charged ions holds them together. Ionic bonds typically form between metals and nonmetals. Metallic Bonds : In metallic bonds, atoms in a metal lattice share a "sea of electrons" that are free to move around. This electron mobility accounts for many physical properties of metals, such as conductivity and malleability. Hydrogen Bonds : A type of weak bond that occurs when a hydrogen atom covalently bonded to a more electronegative atom (like oxygen or nitrogen) is attracted to another electronegative atom. These bonds are crucial in many biological processes, such as the formation of DNA structure and protein folding. A chemical bond is a force that holds atoms together in a molecule or compound. It forms when atoms share or transfer electrons to achieve a more stable electronic configuration, usually resembling the electron configuration of noble gases. There are several types of chemical bonds:

4. Lewis Symbols Valence electrons are reported by dots around the chemical symbol of element, e.g., Ionic Bond A chemical bond formed by complete transference of electrons from one atom (metal) to another (non-metal) and hence, each atom acquire the stable nearest noble gas configuration, is called ionic bond or electrovalent bond, e.g., formation of sodium chloride

5. General properties of ionic compounds: NaCl , Magnesium Oxide ( MgO ), Calcium Carbonate ( CaCO ₃), Potassium Bromide ( KBr ), Lithium Fluoride ( LiF ), Iron Chloride ( FeCl ₂ and FeCl ₃) Physical state:  They form definite pattern that is crystal lattice. Crystal lattice is 3D arrangement of cation and an anion. For example, in NaCl crystal due to crystal formation they all are solids due to strong bonding between constituents. Melting and boiling point:  They have high melting and boiling points because of strong attraction between constituents. Solubility:  We know like dissolves like. So, polar compounds are soluble in polar solvents. Now, ionic compounds have a charge that is they are polar. Therefore, they will dissolve in polar solvents like water. So, all ionic compounds are soluble in polar solvents and insoluble in organic solvents. Electrical conductivity:  It is due to free movement of free ions when ionic compound is dissolved in water. When dissolved they break into ions and conduct electricity.

6. COVALENT BOND They are formed by mutual sharing of electrons between combining atoms Hydrogen chloride molecule Oxygen molecule In Nitrogen molecule

7. Terms to know: Valence electrons: The electrons present in valence shell. Lone pair: The electrons that do not participate in bond formation. Bond pair: The electrons that participate in bond formation. Lewis dot structures: Chlorine molecule Carbondioxide molecule 3. Methane

8. Covalent compounds, formed by atoms sharing electrons, have several distinctive properties. Here are some general characteristics: Low Melting and Boiling Points : Covalent compounds usually have lower melting and boiling points compared to ionic compounds. This is because the forces holding covalent molecules together (van der Waals forces or dipole-dipole interactions) are weaker than the ionic bonds in ionic compounds. Poor Electrical Conductivity : In their solid or liquid states, covalent compounds generally do not conduct electricity well. This is because they lack free-moving charged particles (like ions) that can carry an electric current. However, some covalent compounds can conduct electricity when dissolved in water (e.g., acids) because they may ionize in solution. Solubility : Covalent compounds tend to be soluble in nonpolar solvents (like hexane) but are less soluble in polar solvents (like water). The solubility depends on the principle "like dissolves like"; polar covalent compounds are more likely to dissolve in polar solvents, while nonpolar covalent compounds are more likely to dissolve in nonpolar solvents .

9. Distinctive Colors : Many covalent compounds exhibit a range of colors, especially in organic compounds and transition metal complexes. This is often due to the presence of various functional groups or metal ions that absorb specific wavelengths of light. Molecular Structure : Covalent compounds exist as discrete molecules with specific shapes and bond angles, as determined by the arrangement of atoms and electron pairs around the central atom. The shape and structure influence many of their physical properties. Softness and Flexibility : Many covalent compounds, especially organic ones, are softer and more flexible compared to ionic compounds. For instance, many covalent solids are not as brittle as ionic solids. Low Density : Covalent compounds, particularly those with complex molecular structures, often have lower densities compared to ionic compounds.

10. Coordinate bond Coordinate bond is formed when shared pair of electrons comes only from one atom.  There is no mutual sharing of electrons. The one that donates electron is called donor atom and other is called acceptor. The bond is also called dative bond .

11. Limitations of the Octet Rule 1.The incomplete octet of the central atoms : In some covalent compounds central atom has less than eight electrons, i.e., it has an incomplete octet. For example, Li, Be and B have 1, 2, and 3 valence electrons only.

12.   2.Odd-electron molecules: There are certain molecules which have odd number of electrons the octet rule is not applied for all the atoms. The expanded Octet: In many compounds there are more than eight valence electrons around the central atom. It is termed as expanded octet. For Example,

13.   Other Drawbacks of Octet Theory ( i ) Some noble gases, also combine with oxygen and fluorine to form a number of compounds like XeF 2  , XeOF 2  etc. (ii) This theory does not account for the shape of the molecule. (iii) It does not give any idea about the energy of The molecule and relative stability.   Bond Length It is defined as the equilibrium distance between the centres of the nuclei of the two bonded atoms. It is expressed in terms of A. Experimentally, it can be defined by X-ray diffraction or electron diffraction method.

14. A conjugated bond refers to a system of alternating single and double bonds between atoms in a molecule, typically involving carbon atoms. This concept is most commonly discussed in the context of organic chemistry and is crucial for understanding the behavior of certain molecules, especially in terms of their electronic structure and stability. Key Features of Conjugated Systems Alternating Bonds : In a conjugated system, single and double bonds alternate in a continuous sequence. For example, in the conjugated diene system 1,3-butadiene (CH₂=CH-CH=CH₂), the alternating double bonds are conjugated. Delocalized Electrons : In conjugated systems, the π (pi) electrons from the double bonds are not confined to the individual bonds. Instead, they are delocalized over the entire length of the conjugated system. This delocalization creates a region of electron density that is spread out over several atoms .

15. Stabilization : The delocalization of electrons in conjugated systems often results in increased stability compared to isolated double bonds. This phenomenon is known as resonance stabilization. For example, benzene (C₆H₆) has a conjugated system of alternating double bonds, which is represented by a resonance hybrid of multiple structures, leading to its unusual stability. Color and Absorption : Conjugated systems can absorb light in the visible spectrum, which can impart color to the compound. This is due to the electronic transitions between the delocalized π orbitals. For example, many dyes and pigments contain conjugated systems that are responsible for their vivid colors. Reactivity : Conjugated systems can influence the reactivity of molecules. For instance, conjugated dienes can undergo addition reactions differently than non-conjugated dienes due to their electronic structure.

16. Examples of Conjugated Systems Benzene (C₆H₆) : Benzene has a conjugated system of six carbon atoms with alternating single and double bonds, creating a stable ring structure with delocalized π electrons. 1,3-Butadiene (CH₂=CH-CH=CH₂) : This molecule has a conjugated system of alternating double bonds between four carbon atoms. Carotenoids : These are natural pigments found in plants and fruits with long conjugated systems that absorb visible light and contribute to their color. Polycyclic Aromatic Hydrocarbons (PAHs) : Compounds like anthracene and phenanthrene have extensive conjugated systems across multiple fused aromatic rings. Conjugation is a fundamental concept in organic chemistry, influencing molecular properties, reactivity, and spectroscopy.

17. These systems have several important properties and implications in chemistry: Delocalization of Electrons : In a conjugated system, π-electrons are not confined to a single bond or a single pair of atoms but are delocalized over several adjacent atoms. This delocalization lowers the overall energy of the molecule and can enhance stability, which is known as resonance stabilization. Increased Stability : Conjugated systems are generally more stable than isolated double bonds or non-conjugated systems due to this electron delocalization. This stability is sometimes referred to as "resonance energy." Altered Physical Properties : Color : Conjugated systems can absorb visible light due to their electronic transitions. The extent of conjugation affects the wavelength of light absorbed, which in turn influences the color observed. For example, many dyes and pigments are colored because they contain conjugated systems. Spectroscopy : Conjugated systems show characteristic absorption in UV-Vis spectroscopy. The longer the conjugated system, the longer the wavelength of light absorbed (redshift), which provides information about the extent of conjugation .

18. Chemical Reactivity : Electrophilic Aromatic Substitution : In conjugated aromatic systems, such as benzene and its derivatives, the delocalized π-electrons make the ring more reactive towards electrophiles. Diels-Alder Reaction : Conjugated systems are involved in [4+2] cycloaddition reactions, where a diene (a molecule with two double bonds in conjugation) reacts with a dienophile to form a cyclic compound. Electronic Properties : Conjugated systems can exhibit unique electronic properties, such as in organic semiconductors and conductive polymers. The delocalized π-electrons contribute to their ability to conduct electricity. Molecular Orbitals : The π-electrons in conjugated systems occupy molecular orbitals that are combinations of atomic orbitals. In conjugated systems, these molecular orbitals are spread over the entire conjugated network, influencing the system's electronic properties.

19. Polarity of Bonds Polar and Non-Polar Covalent bonds Non-Polar Covalent bonds: When the atoms joined by covalent bond are the same like; H 2 , 0 2 , Cl 2 , the shared pair of electrons is equally attracted by two atoms and thus the shared electron pair is equidistant to both of them. Alternatively, we can say that it lies exactly in the centre of the bonding atoms. As a result, no poles are developed and the bond is called as non-polar covalent bond. The corresponding molecules are known as non-polar molecules. For Example,

20. Polar bond: When covalent bonds formed between different atoms of different electronegativity, shared electron pair between two atoms gets displaced towards highly electronegative atoms. For Example, in HCl molecule, since electronegativity of chlorine is high as compared to hydrogen thus, electron pair is displaced more towards chlorine atom, thus chlorine will acquire a partial negative charge (δ – ) and hydrogen atom have a partial positive charge (δ + ) with the magnitude of charge same as on chlorination. Such covalent bond is called polar covalent bond.

21. The Valence Shell Electron Pair Repulsion (VSEPR) Theory Sidgwick and Powell in 1940, proposed a simple theory based on repulsive character of electron pairs in the valence shell of the atoms. It was further developed by Nyholm and Gillespie (1957). Main Postulates are the following: ( i ) The exact shape of molecule depends upon the number of electron pairs (bonded or non bonded) around the central atoms. (ii) The electron pairs have a tendency to repel each other since they exist around the central atom and the electron clouds are negatively charged. (iii) Electron pairs try to take such position which can minimize the rupulsion between them. (iv) The valence shell is taken as a sphere with the electron pairs placed at maximum distance. (v) A multiple bond is treated as if it is a single electron pair and the electron pairs which constitute the bond as single pairs .

22. Valence Bond Theory of Covalent Bond According to this theory, a covalent bond is formed by the overlapping of two half-filled atomic orbitals having electrons with opposite spins. It is based on wave nature of electron . Depending upon the type of overlapping, the covalent bonds are of two types, known as sigma (σ ) and pi (π) bonds. 1 .   Sigma Bond  (σ bond ) Sigma bond is formed by the end to end (head-on) overlap of bonding orbitals along the internuclear axis. The following result in the formation of σ bond. ( i ) s-s overlapping (ii) S-p overlapping (iii) P-p head to head overlapping (axial )

23. s-s overlapping:  In this case, there is overlap of two half-filled s-orbitals along the internuclear axis as shown below: s-p overlapping : This type of overlapping occurs between half-filled s-orbitals of one atom and half filled p-orbitals of another atoms.

24. p-p overlapping:  This type of overlapping takes place between half filled p-orbitals of the two approaching atoms.

25. Examples

26. 2.  Pi Bond  (π bond) It is formed by the sidewise or lateral overlapping between p- atomic orbitals [pop side by side or lateral overlapping] π bond is a weaker bond than σ bond. Strength of Sigma and pf Bonds Sigma bond (σ bond) is formed by the axial overlapping of the atomic orbitals while the π-bond is formed by side wise overlapping. Since axial overlapping is greater as compared to side wise. Thus, the sigma bond is said to be stronger bond in comparison to a π-bond. Distinction between sigma and n bonds

27. Distinction between sigma and pi bonds

28. Hybridisation Hybridisation is the process of intermixing of the orbitals of slightly different energies so as to redistribute their energies resulting in the formation of new set of orbitals of equivalent energies and shape. Salient Features of Hybridisation : ( i ) Orbitals with almost equal energy take part in the hybridisation . (ii) Number of hybrid orbitals produced is equal to the number of atomic orbitals mixed, (iii) Geometry of a covalent molecule can be indicated by the type of hybridisation . (iv) The hybrid orbitals are more effective in forming stable bonds than the pure atomic orbitals.

29. Conditions necessary for hybridisation : ( i ) Orbitals of valence shell take part in the hybridisation . (ii) Orbitals involved in hybridisation should have almost equal energy. (iii) Promotion of electron is not necessary condition prior to hybridisation . (iv) In some cases filled orbitals of valence shell also take part in hybridisation .

30. Types of Hybridisation : ( i ) sp hybridisation : When one s and one p-orbital hybridise to form two equivalent orbitals, the orbital is known as sp hybrid orbital, and the type of hybridisation is called sp hybridisation .

31. (ii) sp 2   hybridisation : In this type, one s and two p-orbitals hybridise to form three equivalent sp 2   hybridised orbitals. All the three hybrid orbitals remain in the same plane making an angle of 120°. Example. A few compounds in which sp 2   hybridisation takes place are BF 3 , BH 3 , BCl 3  carbon compounds containing double bond etc.

32.   sp 3   hybridisation : In this type, one s and three p-orbitals in the valence shell of an atom get hybridised to form four equivalent hybrid orbitals. There is 25% s-character and 75% p-character in each sp 3  hybrid orbital. The four sp 3  orbitals are directed towards four corners of the tetrahedron.

33. The angle between sp 3  hybrid orbitals is 109.5°. A compound in which sp 3   hybridisation occurs is, (CH 4 ). The structures of NH 2  and H 2 0 molecules can also be explained with the help of sp 3   hybridisation .

34. In NH3 , the valence shell (outer) electronic configuration of nitrogen in the ground state is sp 3   hybridization in NH 3 molecule

35. In NH3 molecule three unpaired electrons in the sp3 hybrid orbitals and a lone pair of electrons is present in the fourth one. These three hybrid orbitals overlap with 1s orbitals of hydrogen atoms to form three N–H sigma bonds. We know that the force of repulsion between a lone pair and a bond pair is more than the force of repulsion between two bond pairs of electrons. The molecule thus gets distorted and the bond angle is reduced to 107° from 109.5°. The geometry of such a molecule will be pyramidal

36. In case of H2O molecule, the four oxygen orbitals (one 2s and three 2p) undergo sp3 hybridisation forming four sp3 hybrid orbitals out of which two contain one electron each and the other two contain a pair of electrons. These four sp3 hybrid orbitals acquire a tetrahedral geometry, with two corners occupied by hydrogen atoms while the other two by the lone pairs. The bond angle in this case is reduced to 104.5° from 109.5° and the molecule thus acquires a V-shape or angular geometry sp 3   hybridization in H 2 O molecule

37. Therefore in ethane C–C bond length is 154 pm and each C–H bond length is 109 pm

38. sp2 Hybridisation in C 2 H 4 :

39. In the formation of ethene molecule, one of the sp2 hybrid orbitals of carbon atom overlaps axially with sp2 hybridised orbital of another carbon atom to form C–C sigma bond. While the other two sp2 hybrid orbitals of each carbon atom are used for making sp2–s sigma bond with two hydrogen atoms. The unhybridised orbital (2px or 2py ) of one carbon atom overlaps sidewise with the similar orbital of the other carbon atom to form weak π bond, which consists of two equal electron clouds distributed above and below the plane of carbon and hydrogen atoms. Thus, in ethene molecule, the carboncarbon bond consists of one sp2–sp2 sigma bond and one pi (π ) bond between p orbitals which are not used in the hybridisation and are perpendicular to the plane of molecule; the bond length 134 pm. The C–H bond is sp2–s sigma with bond length 108 pm. The H– C–H bond angle is 117.6° while the H–C–C angle is 121°. The formation of sigma and pi bonds in ethene

40. sp Hybridisation in C2H2 : In the formation of ethyne molecule, both the carbon atoms undergo sp -hybridisation having two unhybridised orbital i.e., 2 p y and 2 p x . One sp hybrid orbital of one carbon atom overlaps axially with sp hybrid orbital of the other carbon atom to form C–C sigma bond , while the other hybridised orbital of each carbon atom overlaps axially with the half filled s orbital of hydrogen atoms forming s bonds. Each of the two unhybridised p orbitals of both the carbon atoms overlaps sidewise to form two p bonds between the carbon atoms. So the triple bond between the two carbon atoms is made up of one sigma and two pi bonds as shown in Fig.

41. ( i ) Formation of PCl 5 (sp 3 d hybridisation ): The ground state and the excited state outer electronic configurations of phosphorus (Z=15) are represented below.

42. ( ii) Formation of SF6 (sp3d2 hybridisation ): In SF 6 the central sulphur atom has the ground state outer electronic configuration 3 s 2 3 p 4 . In the exited state the available six orbitals i.e., one s , three p and two d are singly occupied by electrons. These orbitals hybridise to form six new sp 3 d 2 hybrid orbitals, which are projected towards the six corners of a regular octahedron in SF 6 . These six sp 3 d 2 hybrid orbitals overlap with singly occupied orbitals of fluorine atoms to form six S–F sigma bonds. Thus SF 6 molecule has a regular octahedral geometry as shown in Fig

43. Octahedral geometry of SF 6 molecule